View source for Iron(II) carbonate

{{short description|Chemical, compound of iron carbon and oxygen}}
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| ImageFile = Iron(II) carbonate.svg
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| Name = iron(II) carbonate
| OtherNames = ferrous carbonate
|Section1={{Chembox Identifiers
| ChemSpiderID_Ref = {{chemspidercite|??|chemspider}}
| ChemSpiderID = 10774
| ChEMBL_Ref = {{ebicite|??|EBI}}
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| UNII_Ref = {{fdacite|correct|FDA}}
| UNII = MZ3Q72U52O
| InChI = 1S/CH2O3.Fe/c2-1(3)4;/h(H2,2,3,4);/q;+2/p-2
| InChIKey = RAQDACVRFCEPDA-UHFFFAOYSA-L
| SMILES = C(=O)([O-])[O-].[Fe+2]
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| CASNo = 563-71-3 
| CASNo_Ref = {{cascite|correct|CAS}}
| PubChem = 11248
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|Section2={{Chembox Properties
| Formula = FeCO<sub>3</sub>
| MolarMass = 115.854 g/mol
| Appearance = white powder or crystals
| Density = 3.9 g/cm<sup>3</sup><ref name="lide"/>
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| MeltingPt_notes = decomposes
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| Solubility = 0.0067 g/L;<ref name="patty"/> K<sub>sp</sub> = 1.28 × 10<sup>−11</sup> <ref name="weisun"/>
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| SolubilityProduct = 3.13{{e|&minus;11}}<ref name="crc">{{cite book |author1=John Rumble |title=CRC Handbook of Chemistry and Physics |date=June 18, 2018 |publisher=CRC Press |isbn=978-1138561632 |pages=5–188|edition=99 |language=English}}</ref>
| RefractIndex = 
| MagSus = +11,300·10<sup>−6</sup> cm<sup>3</sup>/mol
  }}
|Section3={{Chembox Structure
| Coordination = 6
| CrystalStruct = Hexagonal scalenohedral / Trigonal ({{overline|3}}2/m) <br/>[[Space group]]: R {{overline|3}}c, a = 4.6916 [[angstrom|Å]], c = 15.3796 Å
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|Section4={{Chembox Thermochemistry
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|Section7={{Chembox Hazards
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|Section8={{Chembox Related
| OtherAnions = [[iron(II) sulfate]]
| OtherCations = [[copper(II) carbonate]], [[zinc carbonate]]
  }}
}}
'''Iron(II) carbonate''', or '''ferrous carbonate''', is a chemical compound with formula {{chem|Fe|C|O|3}}, that occurs naturally as the mineral [[siderite]].  At ordinary ambient temperatures, it is a green-brown [[ionic compound|ionic solid]] consisting of [[iron|iron(II)]] [[cations]] {{chem|Fe|2+}} and [[carbonate]] [[anions]] {{chem|C|O|3|2-}}.<ref name="kirk"/>  The compound crystallizes in the same motif as [[calcium carbonate]].  In this motif, the carbonate dianion is nearly planar.  Its three oxygen atoms each bind to two Fe(II) centers, such that the Fe has an octahedral coordination geometry.<ref>{{cite journal |doi=10.1524/zkri.1981.156.3-4.233 |title=Crystal structure refinements of magnesite, calcite, rhodochrosite, siderite, smithonite, and dolomite, with discussion of some aspects of the stereochemistry of calcite type carbonates |journal=Zeitschrift für Kristallographie - Crystalline Materials |date=1981 |volume=156 |issue=3–4 |pages=233–243 }}</ref>

==Preparation==
Ferrous carbonate can be prepared by reacting solution of the two ions, such as [[iron(II) chloride]] and [[sodium carbonate]]:<ref name="kirk"/>

: {{chem|Fe|Cl|2}} + {{chem|Na|2|C|O|3}} → {{chem|Fe|C|O|3}} + 2{{chem|Na|Cl}}

Ferrous carbonate can be prepared also from solutions of an iron(II) salt, such as [[iron(II) perchlorate]], with [[sodium bicarbonate]], releasing [[carbon dioxide]]:<ref name="singer"/>

: {{chem|Fe}}({{chem|Cl|O|4}})<sub>2</sub> + 2{{chem|Na|H|C|O|3}} → {{chem|Fe|C|O|3}} + 2{{chem|Na|Cl|O|4}} + {{chem|C|O|2}} + {{chem|H|2|O}}

Sel and others used this reaction (but with {{chem|Fe|Cl|2}} instead of {{chem|Fe}}({{chem|Cl|O|4}})<sub>2</sub>) at 0.2 [[molarity|M]] to prepare amorphous {{chem|Fe|C|O|3}}.<ref name="sel"/>

Care must be taken to exclude [[oxygen]] {{chem|O|2}} from the solutions, because the {{chem|Fe|2+}} ion is easily [[oxidation|oxidized]] to {{chem|Fe|3+}}, especially at [[acidity|pH]] above 6.0.<ref name="singer"/>

Ferrous carbonate also forms directly on steel or iron surfaces exposed to solutions of carbon dioxide, forming an "iron carbonate" scale:<ref name="weisun"/>

: {{chem|Fe}} + {{chem|C|O|2}} + {{chem|H|2|O}} → {{chem|Fe|C|O|3}} + {{chem|H|2}}

==Properties==

The dependency of the solubility in water with temperature was determined by Wei Sun and others to be
:<math>
  \log K_{\mathit{sp}} = -59.3498 - 0.041377 T - 2.1963/T + 24.5724 \log T  + 2.518 \sqrt{I} - 0.657 I,
</math>
where ''T'' is the absolute temperature in [[kelvin]]s, and ''I'' is the [[ionic strength]] of the liquid.<ref name="weisun"/>

Iron carbonate decomposes at about {{Convert|500-600|C|K}}.<ref>{{Cite journal |date=2010-12-31 |title=Kinetics of Thermal Decomposition of Iron Carbonate |journal=Egyptian Journal of Chemistry |volume=53 |issue=6 |pages=871–884 |doi=10.21608/ejchem.2010.1268 |issn=2357-0245|doi-access=free }}</ref>

==Uses==
Ferrous carbonate has been used as an iron [[dietary supplement]] to treat [[anemia]].<ref name="osol"/> It is noted to have very poor bioavailability in cats and dogs.<ref name="aafco2014">{{cite web|date=2013|title=AAFCO methods for substantiating nutritional adequacy of dog and cat foods (proposed for 2014 publication)|url=https://www.aafco.org/Portals/0/SiteContent/Regulatory/Committees/Pet-Food/Reports/Pet_Food_Report_2013_Midyear-Proposed_Revisions_to_AAFCO_Nutrient_Profiles.pdf|publisher=AAFCO}}</ref>

==Toxicity==

Ferrous carbonate is slightly toxic; the probable oral lethal dose is between 0.5 and 5 g/kg  (between 35 and 350 g for a 70&nbsp;kg person).<ref name="goss"/>

==Iron(III) carbonate==
Unlike iron(II) carbonate, iron(III) carbonate has not been isolated. Attempts to produce iron(III) carbonate by the reaction of aqueous ferric ions and carbonate ions result in the production of [[iron(III) oxide]] with the release of [[carbon dioxide]] or [[bicarbonate]].<ref>{{cite book |author1=Ronald Rich |title=Inorganic Reactions in Water |date=2007 |publisher=Springer Berlin, Heidelberg |isbn=9783540739616 |page=178 |edition=1st |url=https://doi.org/10.1007/978-3-540-73962-3 |language=en |chapter=8 Iron through Hassium|doi=10.1007/978-3-540-73962-3 }}</ref>

==References==

<references>
<ref name="kirk">(1995): "Kirk-Othmer Encyclopedia of Chemical Technology". 4th ed. Volume 1.</ref>
<ref name="patty">Patty, F., ed. (1963): "Industrial Hygiene and Toxicology"; volume II: 'Toxicology". 2nd ed. Interscience. Page 1053.</ref>
<ref name="lide">D R. Lide, ed.(2000): "CRC Handbook of Chemistry and Physics". 81st Edition. Pages 4-65.</ref>
<ref name="osol">
A .Osol and J. E. Hoover and others, eds. (1975): "Remington's Pharmaceutical Sciences". 15th ed. Mack Publishing. Page 775</ref>
<ref name="singer">{{Cite journal |last1=Singer |first1=Philip C. |last2=Stumm |first2=Werner |date=1970 |title=The Solubility of Ferrous Iron in Carbonate-Bearing Waters |url=https://www.jstor.org/stable/41266171 |journal=Journal (American Water Works Association) |volume=62 |issue=3 |pages=198–202 |doi=10.1002/j.1551-8833.1970.tb03888.x |jstor=41266171 |issn=0003-150X}}</ref>
<ref name="goss">Gosselin, R.E., H.C. Hodge, R.P. Smith, and M.N. Gleason. Clinical Toxicology of Commercial Products. 4th ed. Baltimore: Williams and Wilkins, 1976., p. II-97</ref>
<ref name="weisun">Wei Sun (2009): "Kinetics of iron carbonate and iron sulfide scale formation in CO2/H2S corrosion".  PhD Thesis, Ohio University.</ref>
<ref name="sel">Ozlem Sel, A.V. Radha, Knud Dideriksen, and Alexandra Navrotsky (2012): "Amorphous iron (II) carbonate: Crystallization energetics and comparison to other carbonate minerals related to CO2 sequestration". Geochimica et Cosmochimica Acta, volume 87, issue 15, pages 61–68. {{doi|10.1016/j.gca.2012.03.011}}</ref>
</references>

{{carbonates}}
{{iron compounds}}
{{Antianemic preparations}}

[[Category:Iron(II) compounds]]
[[Category:Carbonates]]